### The Chemical Reaction

Since the dawn of our existence, we humans have been conducting one particular type of chemical reaction and have derived from it both light and energy.

Chemists call this reaction combustion, and whether we are talking about burning logs, coal, or oil, this term refers broadly to any reaction in which a substance combines with oxygen to form carbon dioxide and water, as shown in the equation below:

$\ce{(Substance) + O2 -> CO2 + H2O}.$

In this exploration, we will use combustion to establish how we can compare the energy of different molecules and what that comparison means for us.

# Hiking the Reaction Landscape

When we think of common fuels, one that may come to mind is octane. Octane, pictured below, is a chemical component of gasoline, but many people also use its name to refer to anything that is powerful, intense, or high-energy as in the phrase "high-octane." Why might octane have come to signify a high-energy material?

# Hiking the Reaction Landscape

All it takes is a spark and oxygen for octane to release energy as heat and light.

That energy has to come from somewhere, and during combustion it is actually derived from the particular arrangement of atoms in octane's molecular structure, much as energy is stored in a battery. Oxygen molecules contain energy, too, and it is the combination of octane and oxygen that allows — indeed, kick-starts — the release of energy.

Neither octane nor oxygen on their own can accomplish this.

However, we can’t really get much energy out of $\ce{CO2}$ in this way — it can't release significant energy through combustion. In fact, as the general equation for combustion below shows, it is effectively the end of the road:

$\ce{(Substance) + O2 -> CO2 + H2O}.$

Carbon dioxide is always a product, and never a reactant, in a combustion reaction.

# Hiking the Reaction Landscape

We can visualize these molecules and their relative energy content if we think of them positioned somewhere on an "energy landscape," where hills represent high energy and valleys represent low energy.

When a reaction occurs, it tends to consume its reactants and form products that are energetically "downhill." Just as a ball at the very top of a hill will roll downwards and lower its gravitational potential energy given a sufficient nudge, a molecule like octane will likewise tend to react in a way that lowers its energy, in the end forming lower-energy products like $\ce{ CO2 }$ and $\ce{ H2O }$.

# Hiking the Reaction Landscape

The higher the starting material is on the landscape relative to the valley where its products dwell, the more energy it will release.

# Hiking the Reaction Landscape

For a combustion reaction, regardless of what molecule we start with, the products are always the same: $\ce{ CO2 }$ and $\ce{ H2O }$. This means that, on the energy landscape, all the starting materials that undergo this type of reaction are located on hills surrounding a central valley containing those products.

Out of the starting materials depicted above, which would release the most energy in a combustion reaction?

# Hiking the Reaction Landscape

One thing that bears mentioning is that, for every reaction, there is a trail (sometimes even multiple trails) that leads from materials atop a hill downward into the valley of products. And just like stations along a climbing trail, there are other structures dotting the trail all the way down.

As it traverses the landscape from starting material to product, the reaction makes temporary stops at these points, but it ultimately continues on its way to the valley of final products. Chemists refer to these temporary stops as intermediates — these are not the final products, but they do form transiently, can be observed, and dot the landscape of nearly every reaction pathway.

This analogy will suffice for now, but later we'll see that intermediates are often higher in energy than the starting materials themselves, making for quite a hike to get down into the valley of products!

# Hiking the Reaction Landscape

At this point, it's logical to wonder whether carbon dioxide is just a boring old product to be ignored rather than used for something else. What we can do with $\ce{ CO2 }$ chemically lies, in reality, in the opposite direction.

It turns out that, under certain conditions, $\ce{ CO2 }$ molecules can react with other molecules to build up new, larger ones. For example, during the process known as photosynthesis, plants take energy from the sun and use $\ce{ CO2 }$ as a building block to make molecules of sugar $(\ce{C6H12O2}).$ The equation for this process looks a bit like the one for combustion, only in reverse:

Given this information, would we say that a sugar molecule is higher or lower in energy content than a $\ce{ CO2 }$ molecule?

# Hiking the Reaction Landscape

At this point, we can say that sugar molecules, like octane molecules, lie somewhere uphill from $\ce{ CO2 }$ on our energy landscape. To build molecules from $\ce{ CO2 }$ necessarily requires that we input energy, or "go uphill."

# Hiking the Reaction Landscape

Sugar is not as useful as an everyday fuel as octane, but it still contains a good deal of useful energy, certainly more than $\ce{ CO2 }$. In fact, our bodies use sugar, too! The primary reaction going on inside our cells involving the sugar we consume is called cellular respiration, and the overall reaction can be represented by the equation below:

$\ce{C6H12O6 + 6O2 → 6CO2 + 6H2O}.$

Given what we’ve discussed so far and the equation above, is this reaction more likely to require an investment of energy or release energy overall?

Note: The numbers in front of the chemical symbols are called coefficients, and serve only for bookkeeping; we'll discuss these more explicitly in a later chapter.

# Hiking the Reaction Landscape

Although humans have been reaping the energetic benefits of combustion for centuries, understanding that molecules lie at different points on the energy landscape has become crucial to chemists and chemical engineers. Seeing where materials lie on this energy landscape allows us to make some very powerful predictions about how they can be expected to behave under various conditions, or how they might interact with one another.

It's important to realize, however, that different landscapes exist for different reactions, and a single molecule like octane can occupy vastly different points on these landscapes. It's easy to imagine plenty of landscapes wherein octane might lie in a valley and $\ce{ CO2 }$ might dwell atop a peak!

# Hiking the Reaction Landscape

Lest we get the false impression that $\ce{ CO2 }$ is truly a chemical dead-end, let's imagine a totally different reaction landscape.

This time, instead of kick-starting octane combustion with the energy of $\ce{O2},$ let's imagine $\ce{ CO2 }$ atop a hill, now requiring the energy of the reactive gas called fluorine $(\ce{ F2 })$ to start a reaction.

Notice how the tables are now turned! Just by changing the conditions, we can take $\ce{ CO2 }$ and transform it into a completely different molecule in a downhill process, ultimately also forming $\ce{ O2 }$ as a product.

Chemistry requires us to specify the context — that is, the reaction type and conditions — before we can have a meaningful discussion of relative energies. Substances occupying a low valley in one landscape could easily lie atop a high mountain in another.

# Hiking the Reaction Landscape

Although we’ve only scratched the surface of its potential, we might say that chemistry is ultimately the art of rearranging molecules to accomplish a goal, whether this goal is to construct something large and complex, or instead to extract energy from large molecules to drive another process.

Throughout this course, we’ll explore how reactions work on the molecular scale, and eventually how the properties of atoms influence the structure and behavior of larger molecules. We’ll even begin to understand how nature harnesses basic chemical principles to accomplish reactions of dizzying specificity and complexity within our own bodies!

# Hiking the Reaction Landscape

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